organizing the periodic table - mrs. gross's...
TRANSCRIPT
Organizing the Periodic Table
How did chemists begin to organize the
known elements?
Chemists used the properties of the
elements to sort them into groups.
The Organizers
JW Dobereiner
grouped the elements into triads.
Problem: Not all of the elements could
be put into triads.
The Organizers
Dmitri Mendeleev
arranged the elements by
increasing atomic mass.
Problem: the chemical properties of
the elements did not line up correctly.
Pg. 156
Figure
6.3
The Organizers
Lother Meyer
also created a periodic table similar to
Mendeleev’s, however Mendeleev
published his table first and could
explain his table better.
Development of Atomic Theory
The structure of the atom was
determined through years of
experimentation. Once the structure
of the atom was determined by Dalton,
Thomson, Rutherford, Chadwick, Bohr,
and others, Henry Moseley
determined the atomic number for
the elements.
Organization of
Modern Periodic Table
The elements are arranged
in order of
increasing atomic number.
Periodic Law
When elements are arranged in
order of increasing atomic
number, there is a periodic
repetition of their physical and
chemical properties.
Three classes of elements:
Metals
Nonmetals
Metalloids
Across a period, the properties of
elements become less metallic and
more nonmetallic.
Properties of Metals
•Solids (except for Hg) at room temperature
•Make up 80% of the periodic table
•Good conductors of heat and electricity
•Shiny
•Ductile (can be made into thin wire)
•Malleable (can be hammered into sheets)
Properties of Nonmetals
Properties vary among the nonmetals.
At room temperature, some are gases,
some are solids, and 1 is a liquid.
•Poor conductors
•Brittle
•Not ductile
•Not malleable
Metalloids (Semi-metals)
B, Si, Ge, As, Sb, Te, At
Behave like metals under certain
conditions.
Behave like nonmetals under
different conditions.
Which of these sets of elements have
similar physical and chemical properties?
a. oxygen, nitrogen, carbon, boron
b. strontium, magnesium, calcium, beryllium
c. nitrogen, neon, nickel, niobium
Identify each element as a metal,
metalloid, or nonmetal.
a. gold
b. silicon
c. sulfur
d. barium
Name two elements that have
properties similar to those of the
element sodium.
What pattern is revealed when the
elements are arranged in a periodic table
in order of increasing atomic number?
Identify each property below as more
characteristic of a metal or a nonmetal.
a. a gas at room temperature
b. brittle
c. malleable
d. poor conductor of electric current
e. shiny
In general, how are metalloids different
from metals and nonmetals?
Representative Elements
Elements in columns
1, 2, 13, 14, 15, 16, 17, 18
OR
Elements in columns
1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A
OR
Elements in blocks s and p
Representative Elements
Called “REPRESENTATIVE”
elements because they display a
wide range of chemical and
physical properties.
Which of the following are symbols for
representative elements:
Na, Mg, Fe, Ni, Cl ?
Transition Metals
Elements in columns
3, 4, 5, 6, 7, 8, 9, 10, 11, 12
OR
Elements in columns
3B, 4B, 5B, 6B, 7B, 8B, 1B, 2B
OR
Elements in the d block
Transition Metals
Characterized by electrons
filling the d orbitals
When transition metal ions form compounds,
the compounds often exhibit various colors.
Iron oxide Red
Cadmium sulfate Orange
Copper(II) chloride Blue-green
Inner Transition Metals (Rare earth elements)
Elements in the f block
Consists of the
Lanthanide (La) series
and the
Actinide (Ac) series.
Inner Transition Metals (Rare earth elements)
Contains some elements that occur
naturally
and some that are synthetic.
Some are radioactive, some are not.
Groups in the Periodic Table
Alkali metals – column 1 (or 1A)
Alkaline earth metals – column 2 (or 2A)
Halogens (salt-formers) – column 17 (7A)
Noble gases – column 18 (or 8A)
Why do the elements potassium and
sodium have similar chemical properties?
Classify each element as a
representative element, transition
metal, or noble gas.
a. 1s22s22p63s23p64s23d104p6
b. 1s22s22p63s23p64s23d6
c. 1s22s22p63s23p2
Which of the following elements are
transition metals:
Cu, Sr, Cd, Au, Al, Ge, Co ?
How many electrons are in the
highest occupied energy level of a
Group 5A element?
Trends
in the
Periodic
Table
Shielding is the result of the core
electrons blocking the ‘pull’ of the
nucleus on the outermost electrons.
Terms to Know
Nuclear charge is the charge
in the nucleus.
Terms to Know
Atomic Radius is one half the distance
between the nuclei of two atoms of the
same element when the atoms are joined.
Pg. 170 Figure 6.13
Atomic Radius
From left to right across a period,
the atomic radius generally decreases.
Why? As you go across a row, there is an increase
in the number of protons and the number of electrons.
However, the electrons being added are added to the
same principle energy level, so the shielding remains
the same. Therefore, the increased nuclear charge
pulls the outermost electrons in closer to the nucleus.
The result is a smaller atom.
(See pg. 171, 3rd paragraph.)
Atomic Radius
From top to bottom down a group, the
atomic radius generally increases.
Why? As you go down a column, there is an
increase in the number of protons and the number of
occupied energy levels. The increase in positive
charge pulls the electrons closer to the nucleus, but
the increase in the number of occupied orbitals
shields the nucleus more. The shielding effect is
greater than the increased nuclear charge. The
result is a larger atom.
(See pg. 171, 1st and 2nd paragraph.)
Which element in each pair has atoms
with a larger atomic radius?
a. sodium, lithium
b. strontium, magnesium
c. carbon, germanium
d. selenium, oxygen
Explain why fluorine has a smaller atomic
radius than both oxygen and chlorine.
Terms to Know
Ion is a neutral atom that has either lost
or gained an electron.
Atoms that have gained electrons are
called
Atoms that have lost electron(s) are called
Positive Ions or Cations
Negative Ions or Anions
Cation is an ion that has lost one or
more electrons resulting in a net
positive charge.
Anion is an ion that has gained one
or more electrons resulting in a net
negative charge.
Ion Size
Cations are always
SMALLER than the atoms
from which they form.
Why? When a metal atom loses an electron,
the attraction between the remaining electrons
and the nucleus is increased. The electrons are
drawn closer to the nucleus.
(See pg. 176, 2nd paragraph)
Ion Size
Anions are always
LARGER than the atoms
from which they form.
Why? As the number of electrons
increases, the attraction of the nucleus
for any one electron decreases.
(See pg. 176, 3rd paragraph)
Li F O N C B Be
Li1+ F1- O2- N3- C B3+ Be2+
Ion Trends
From left to right across a period, the
cation size decreases.
From left to right across a period, the
anion size decreases.
(See figure 6.20 on pg. 176)
Pg. 176 Figure 6.20
From top to bottom down a group,
the ionic size increases.
(Pg. 172)
Ion Trends
Which particle has the larger
radius in each atom/ion pair?
a. Na, Na+
b. S, S2-
c. I, I-
d. Al, Al3+
In each pair, which ion is larger?
a. Ca2+, Mg2+
b. Cl - , P3-
c. Cu+, Cu2+
The atomic radius _____________ from
left to right across the periodic table.
The atomic radius ____________
from top to bottom on the period table.
Aside from the noble gases, what is the
smallest atom on the periodic table?
Aside from the noble gases, what is the
largest atom on the periodic table?
The atom is _________ than its cation.
The atom is _________ than its anion.
Cations are _________ than the
anions that are in the same row.
The cation trend from left to right is
that the ionic radius ____________.
The anion trend from left to right is
that the ionic radius _____________.
List the following from smallest to
largest according atomic radius.
a. Mg, S, Na, Si
b. Be, Ba, Ca, Ra
c. F, As, Br, Ga, Cl
List the following from smallest to
largest according ionic radius.
a. Mg2+, S2-, Na+, Cl-
b. Be2+, Ba2+, Ca2+, Ra2+
c. F-, As3-, Br-, Ga3+, Cl-
Terms to Know
Ionization
Energy
is the energy
required to
remove an
electron
from an
atom. (Pg. 173-175)
Explain the difference between the terms
“first ionization energy” and “second
ionization energy” of an element.
The first ionization energy is the energy
needed to remove a first electron from an
atom.
The second ionization energy is the energy
needed to remove a second electron.
Why is there a large increase between the
first and second ionization energies of the
alkali metals?
It is relatively easy to remove the first
electron from an alkali metal atom because
they have only one valence electron;
it is much more difficult to remove the
second because it is being removed from a
full energy level.
There is a large jump between the
second and third ionization energies
of magnesium. There is a large jump
between the third and fourth
ionization energies of aluminum.
Explain these observations.
It is relatively easy to remove two electrons
from magnesium because it has 2 valence
electrons; it is much more difficult to remove a
third electron because it comes from a full
energy level.
It is relatively easy to remove three electrons
from aluminum because it has 3 valence
electrons; it is much more difficult to remove a
fourth electron because it comes from a full
energy level.
Ionization Energy
From left to right across a period,
the ionization energy generally
increases.
Why? As you go across a row, there is an increase
in the nuclear charge, but the shielding effect remains
the same. There is an increase in the attraction of the
nucleus for an electron. Thus it takes more energy to
remove an electron from the atom.
(See pg. 174, 2nd paragraph.)
From top to bottom down a group,
the ionization energy generally
decreases.
Why? As the size of the atom increases and the
amount of shielding electrons increases, nuclear
charge has a smaller effect on the electrons in the
outermost energy level. Therefore, less energy is
required to remove an electron from the outermost
energy level.
(See pg. 174, 1st paragraph.)
Ionization Energy
Which element in each pair has a
greater first ionization energy?
a. lithium, boron
b. magnesium, strontium
c. cesium, aluminum
Would you expect metals or nonmetals in
the same period to have higher ionization
energies? Give a reason for your answer.
Explain each of the following comparisons.
Calcium has a smaller second ionization
energy than does potassium.
Explain each of the following comparison.
Lithium has a larger first ionization
energy than does cesium.
Explain the following comparison.
Magnesium has a larger third
ionization energy than does aluminum.
Electronegativity
Electronegativity
is the ability of an atom to attract
electrons when the atom is in a
compound.
(See pg. 177, 1st paragraph)
For the representative elements:
From top to bottom down a group,
the electronegativity values decrease.
(See pg. 177, 3rd paragraph)
From left to right across a period,
the electronegativity values increase.
Electronegativity
Electronegativity
Fluorine is the most electronegative
element on the periodic table.
Why are the noble gases NOT included
in electronegativity table?
Because they are UNREACTIVE and
do NOT want to bond with any atoms.
Which element in each pair has a
higher electronegativity value?
a. Cl, F
b. C, N
c. Mg, Ne
d. As, Ca
When the elements in each pair
are chemically combined, which
element in each pair has a
greater attraction for electrons?
a. Ca or O
b. O or F
c. H or O
d. K or S
Pg. 171
Figure 6.14
Pg. 174
Figure 6.17